Draeger
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Copper bicarbonate?
So, yesterday I decided to use sodium bicarbonate to convert some copper waste into basic copper carbonate. What ended up happening after the bubbling
stopped is it forming a grey solution of unknown composition. What could it be?
Collected elements:
Al, Cu, Ga, C (coal), S, Zn, Na
Collected compounds:
Inorganic:
NaOH; NaHCO3; MnCl2; MnCO3; CuSO4; FeSO4; aq. 30-33% HCl; aq. NaClO; aq. 9,5% ammonia; aq. 94-96% H2SO4; aq. 3% H2O2
Organic:
citric acid, sodium acetate, sodium citrate, petroleum, mineral oil
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Texium
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Not "copper bicarbonate" that's for sure. What else was in your waste solution? Since you mentioned bubbling, I presume there must have been some kind
of acid present. Do you have litmus or universal indicator strips that you could use to check the pH of the solution?
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woelen
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Copper bicarbonate does not exist. When you mix copper(II) ions and bicarbonate in solution, you get copper carbonate, CO2 and water. I even doubt
that you get pure copper carbonate, most likely it is somewhat basic copper carbonate, also containing hydroxide.
The copper carbonate you get has a blue/cyan color. If your material is grey, then probably it contains other metal cations as well. You write about
'copper waste', so I assume your material is not a pure copper salt.
If also chloride is present in the solution, then things get more complicated. Chloride ions keep coordinated to the copper metal ions as well. In
that case, you'll get a complicated copper carbonate/chloride/hydroxide mix. This mix has a nice green color (provided, you only have copper metal
cations in your mix).
A nice experiment is making a precipitate of a solution of copper sulfate with sodium hydroxide, and another with sodium (bi)carbonate. Then repeat
the experiment with a solution of copper sulfate, to which some table salt is added as well. You'll see that you get completely differently colored
precipitates, due to the coordinated chloride ions.
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Draeger
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Quote: Originally posted by woelen  | Copper bicarbonate does not exist. When you mix copper(II) ions and bicarbonate in solution, you get copper carbonate, CO2 and water. I even doubt
that you get pure copper carbonate, most likely it is somewhat basic copper carbonate, also containing hydroxide.
The copper carbonate you get has a blue/cyan color. If your material is grey, then probably it contains other metal cations as well. You write about
'copper waste', so I assume your material is not a pure copper salt.
If also chloride is present in the solution, then things get more complicated. Chloride ions keep coordinated to the copper metal ions as well. In
that case, you'll get a complicated copper carbonate/chloride/hydroxide mix. This mix has a nice green color (provided, you only have copper metal
cations in your mix).
A nice experiment is making a precipitate of a solution of copper sulfate with sodium hydroxide, and another with sodium (bi)carbonate. Then repeat
the experiment with a solution of copper sulfate, to which some table salt is added as well. You'll see that you get completely differently colored
precipitates, due to the coordinated chloride ions. |
Thank you. Yeah, I was just stupid. I wrote it at such a late hour that I just forgot that I label my waste bags after what hazardous substance is in
it, not generally what is in it. It's probably contaminated with lots of sodium and chloride, and even probably lots of calcium.
Collected elements:
Al, Cu, Ga, C (coal), S, Zn, Na
Collected compounds:
Inorganic:
NaOH; NaHCO3; MnCl2; MnCO3; CuSO4; FeSO4; aq. 30-33% HCl; aq. NaClO; aq. 9,5% ammonia; aq. 94-96% H2SO4; aq. 3% H2O2
Organic:
citric acid, sodium acetate, sodium citrate, petroleum, mineral oil
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Syn the Sizer
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| Quote: Originally posted by woelen | If also chloride is present in the solution, then things get more complicated. Chloride ions keep coordinated to the copper metal ions as well. In
that case, you'll get a complicated copper carbonate/chloride/hydroxide mix. This mix has a nice green color (provided, you only have copper metal
cations in your mix).
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If your copper waste has chloride ions in it, is it best just to crash the copper out with aluminum? I realize you will end up with copper metal
instead of a easily usable basic carbonate. Is it possible to prevent the ion coordination?
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chornedsnorkack
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| Quote: Originally posted by woelen | Copper bicarbonate does not exist. When you mix copper(II) ions and bicarbonate in solution, you get copper carbonate, CO2 and water. I even doubt
that you get pure copper carbonate, most likely it is somewhat basic copper carbonate, also containing hydroxide.
The copper carbonate you get has a blue/cyan color. |
Pure copper carbonate stability field in room temperature water was found to be above 4,6 bar carbon dioxide. And judging by repeated failed attempts
to produce it, probably not kinetically straightforward.
Copper has two basic carbonates - azurite Cu3(CO3)2(OH)2 and malachite
Cu2CO3(OH)2. Different colours - one blue, one green. How do you reproducibly produce one or other?
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Bezaleel
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In case you are looking to purify you copper carbonate, Atomistry suggests to dissolve the hydroxide in ammonia solution, which leaves pure CuCO3
behind.
I think this will also help you get rid you the chloride, forming dissolved NH4Cl.
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Bedlasky
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| Quote: Originally posted by Bezaleel | In case you are looking to purify you copper carbonate, Atomistry suggests to dissolve the hydroxide in ammonia solution, which leaves pure CuCO3
behind.
I think this will also help you get rid you the chloride, forming dissolved NH4Cl. |
Ammonia solution also dissolve copper carbonate.
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